Properties of Solids

• Kinetic molecular theory provides an overview of the microscopic properties of molecules or atoms and their interactions.
• The kinetic molecular theory describes the microscopic properties of matter and how they translate to states of matter.
• Matter is composed of small particles.
• Molecules interact through attractive forces.
• Molecules are in constant, random motion.
• The amount of kinetic energy in a substance is related to its temperature.

States of Matter

• Solid: Has a fixed shape and volume, is incompressible, has high density, cannot flow, has high rigidity, particles are tightly packed, and has low kinetic energy with strong attraction between particles.
• Liquid: Takes the shape of the container, has a fixed volume, is only slightly compressible, has low density, can flow, has less rigidity, particles are loosely packed, and has moderate kinetic energy with moderate attraction between particles.

Intermolecular Forces of Attraction (IMFA)

• Intermolecular forces are attractive forces between molecules and are weaker than intramolecular forces.
• Dipole-dipole, dipole-induced dipole, and dispersion forces are van der Waals forces.

Dipole-Dipole Forces

• Attractive forces between polar molecules.
• The positive side of a polar molecule attracts the negative side of another polar molecule.
• The larger the dipole moment, the greater the force.
• Examples of molecules exhibiting dipole-dipole forces: HCl and H2S.

Hydrogen Bonding

• A special type of dipole-dipole force.
• An attractive force that exists when hydrogen is bonded to the most electronegative atoms, namely fluorine, oxygen, or nitrogen.
• Examples of molecules exhibiting hydrogen bonding: HF, H2O, NH3

Ion-Dipole Forces

• Attractive forces between an ion and a polar molecule.
• A cation interacts with the partially negative side of a polar molecule, while an anion interacts with a partially positive side.
• Coulomb's law is used to determine which ion will have the strongest or weakest interactions with the polar molecule.
• The force's magnitude is directly related to the charge magnitude on the particles and indirectly related to the distance between the particles.
• The strength of this IMFA kind increases as the ion charge increases, while the distance decreases as the force strength increases.
• Example: NaCl (aq).

London Dispersion Forces

• Named after Fritz London.
• Present between all electrically neutral molecules.
• London dispersion forces are the weakest type of IMFA.
• Arise from temporary dipoles (instantaneous dipoles) induced in atoms or molecules.
• Polarizability refers to the ease at which the electron cloud can be distorted.
• Examples: He, H2.

Ion-Induced Dipole Forces

• Occur when an ion interacts with a nonpolar molecule.
• The charge on the ion induces polarity in the nonpolar molecule.
• Example: Fe2+ & O2.

Dipole-Induced Dipole Forces

• Weak attraction when a polar molecule induces a dipole in an atom or nonpolar molecule by disturbing the arrangement of electrons.
• Examples: HCl & Ar, H2O & O2.

Properties of Liquids

Surface Tension

• The tendency of a fluid to resist an external force and acquire the least possible surface area.
• Molecules in the interior are attracted equally on all sides, while those at the surface are attracted only below and to the sides, thus producing a net inward force.
• Surface tension is higher in liquids with higher intermolecular forces.

Viscosity

• The measure of a fluid's resistance to flow.
• Liquids that flow easily have low viscosity, while liquids that do not flow readily have high viscosity.
• Liquids with stronger IMFA have a higher viscosity than those with weaker IMFA.
• The more massive the liquid, the higher the viscosity.
• Molecules with long chains have a higher viscosity than shorter molecules.

Solubility

• The ability of a substance to dissolve in a given amount of solvent at a specified temperature.
• Solute and solvent with similar IMFA form a solution.
• Polar molecules are miscible (mixable)in a polar solvent.
• Nonpolar molecules are miscible in a nonpolar solvent.
• Polar and nonpolar molecules do not mix together.
• Stronger intermolecular forces between the solvent and solute result in greater solubility.

Vapor Pressure

• The pressure exerted by a vapor in equilibrium with its liquid phase in a closed system.
• Occurs when the rate of evaporation equals the rate of condensation.
• Substances with stronger IMFAs have lower vapor pressure compared to those with weaker IMFA.

Heat of Vaporization (ΔΗvap)

• The energy required to transform a quantity of liquid into a gas. Also known as "enthalpy of vaporization" or "molar heat of vaporization." –Vaporization occurs more readily with increased temperature, increased surface area of the liquid, and decreased strength of intermolecular forces
• Stronger intermolecular forces result in a higher heat of vaporization.

Boiling Point

• The temperature at which a substance changes from liquid to gas.
• Boiling happens when the molecules of a liquid gain enough energy to overcome the intermolecular forces of attraction.
• Liquids with stronger intermolecular forces tend to have higher boiling points.
• Larger molecules also have more surface area for intermolecular interactions, which means more energy is required to break these interactions during boiling.

Properties of Solids

• Crystalline solids are solids with a definite repeating pattern in which atoms, ions, or molecules are arranged in.
• A lattice is the orderly arrangement of atoms in crystals.
• Lattice point is a specific atom, molecule, or ion.
• Crystalline solids are formed from repeating units called unit cells.
• Unit cells are repeated, giving rise to a 3-D crystal structure known as a crystal lattice.
• Covalent solids are bonded by atoms which are covalently bonded to one another. Examples include diamond, graphite, charcoal, silicon dioxide/ quartz.
• Metallic solids are composed of metal atoms bonded together by metallic bonds. Metallic bonds have positive metal ions that are attracted by a “sea of electrons.”
• Ionic Crystals/ Ionic Solids are made up of alternating positive and negative ions held together by strong coulombic forces.
• Molecular solids are held together by weak intermolecular forces.

Amorphous Solids

• Lacks the order found in crystilline solids.
• Structures at the atomic level are similar to the structures of liquids.
• Atoms, ions, or molecules have little freedom to move, unlike in liquids.
• Don’t have well-defined shapes of a crystal.

Crystalline vs Amorphous solids

• Crystalline solids are held together by uniform, strong intermolecular forces of attraction.
• Amorphous solids are held together by non-uniform intermolecular forces of attraction.

Melting/Freezing Point

• Crystalline solids have a precise melting point.
• Amorphous solids melt over a wide range of temperature.
• Heat of fusion (ΔΗfus) is the quantity of heat necessary to melt a solid.
• Crystalline solids have a fixed and define heat of fusion.
• Amorphous solids have no precise heat of fusion value.

Heat of Sublimation (ΔΗsub)

– Quantity of heat necessary to sublime a solid

• Crystalline solids are fixed and definite.
• Amorphous solids have no precise value.

Chemical Nature

• Crystalline solids exhibit anisotropy, where physical and mechanical properties vary with direction.
• Amorphous solids exhibit isotropy, where physical and mechanical properties are equal in all directions.

Other Properties of Solids

• Malleability: the ability of a solid to undergo compressive stress without breaking it.
• Ductility: the ability of solid to undergo tensile stress.
• Electrical Conductivity: the measurement of the ability of atoms, molecules, or ions to transfer electrons from one to another.
• Thermal Conductivity: the measurement of the ability of atoms, molecules, or ions to move and collide with neighboring particles.
• A phase of matter refers to the physical form of matter (solid, liquid, gas, plasma) based on temperature and pressure.
• A state of matter refers to the chemical composition of matter (element, compound, mixture).
• The same substance can exist in different phases of matter, but its state of matter remains constant.
• Phase changes occur when matter changes from one physical form to another (e.g., melting, boiling, condensing), while state changes occur when matter undergoes a chemical reaction (e.g., rusting, burning).

Phase Changes Between States of Matter

• Melting: Solid to liquid.
• Freezing: Liquid to solid.
• Vaporization: Liquid to gas.
• Condensation: Gas to liquid.
• Deposition: Gas to solid.
• Sublimation: Solid to gas.

Phase Diagram

• Represents the various physical states or phases of matter at different pressures and temperatures.
• Summarizes the effects of pressure and temperature on the nature of a substance.
• Consider the temperature as the X-axis and pressure as the Y-axis.
• Using coordinates (°C, atm), plot points and then check where it’s located.
• The A-Bis where the solid phase is in equilibrium with its gaseous counterpart.
• The point where the lines intersect is called the triple point. It is the pressure and temperature conditions at which all three phases, solid, liquid, and gas, coexist in equilibrium.
• The critical point is the highest temperature and pressure at which a gas and a liquid can coexist at equilibrium.
• Beyond the critical point, the two phases become indistinguishable and form a phase known as a supercritical fluid.

Heating and Cooling Curve

• In an endothermic reaction, energy is absorbed in the form of heat from its surroundings, whereas an exothermic reaction releases energy to the surroundings.
• At a sloped line, temperature changes as heat is added.
  • Q=mCpΔT or Q=mCp(Tfinal - Tinitial)
    • Where:
      • Q = heat energy
      • m = the mass (in grams) of the substance
      • C the specific heat capacity
      • ΔT is the change in temperature
• At the straight horizontal line: Heat is added, but temperature remains constant
  • Q=m×H
    • Where:
      • Q is the heat energy
      • m is the mass (in grams) of the substance
      • H is the enthalpy of phase changes (e.g., heat of fusion, vaporization, etc.)
• Calculate the total heat added by solving for the summation of heat per line. (Q1+Q2+Q3+Q4+...)

Concentration of Solution

• Concentration refers to how much solute is present in a solution.

• It is an intensive property.

• Percentage by Mass % m/m: also known as mass percent or weight percent (% w/w).

• Defined as the mass of solute per mass of the solution.

  • The masses of the solute and the solution should be expressed in the same units
  • The mass fraction is unitless

• Percentage by Volume, % v/v: Also known as volume percent (% v/v), is defined as the volume of solute per volume of the solution.

  • Proof is a commonly used unit for alcohol concentration
  • Mathematically defined as two times the percentage by volume of the alcohol in the solution.

• Percentage by Mass per Volume: Also known as the percentage by weight per volume (% w/v), is defined as the mass of solute per volume of the solution.

• The mass of the solute is usually expressed in grams, while the volume of the solution is usually expressed in milliliters.

• Molarity and Molality of a SolutionMolarity (M)

  • Usually expression of concentration of aqueous solution
  • Defined as the number of moles of the solute dissolved in one liter of the solution
  • Easy to use especially for aqueous solutions where the solute is solid.
  • Susceptible to temperature changes (Temperature dependent)

M = nsolutevolume = mol soluteL solution

• Molality (m)
  • m = mol solute/kg solvent
  • Temperature independent
  • Unit: mol/kg (molal) colligative properties

Colligative Properties

• Properties of solutions that depend on the concentration of solutes instead of the identity(nature) or mass of solute particles3 colligative properties exist, including:

  • Vapor Pressure
  • Boiling Point Elevation
  • Freezing Point Depression

• Vapor Pressure

  • The measure of the tendency of a material to change into a gaseous/vapor stateIncreases with temperature
  • Raoult's Law

• Boiling Point Elevation Phenomenon where a solvent's boiling point increases due to the addition of a solute Higher concentration of solute equates to a higher boiling point

• Tb = Tb - Tb° (BP of solution - BP of solvent)

  • (For non electrolyte) - Tb = Kbm ((BPE constant)(molal concentration))
  • (For electrolyte) - Tb = iKbm(Same as previous but with van't Hoff factor)

Freezing Point Depression

• Phenomenon where a solvent's freezing point decreases due to the addition of solute
• Higher concentration equates to a lower freezing point
• Tf = Tf - Tf° (FP of solution - FP of solvent)(For non-electrolyte) - Tb = Kfm
• (For electrolyte) - Tf = iKfm

Test your knowledge of solid-state chemistry. This quiz covers types of solids, their properties (like ductility and conductivity), phase transitions, and the relationship between temperature and kinetic energy. Explore crystalline vs. amorphous solids and intermolecular forces.